Class 11 Chemistry Notes Chapter 3 (Chapter 3) – Examplar Problems (English) Book

Detailed Notes with MCQs of Chapter 3: Classification of Elements and Periodicity in Properties from your NCERT Exemplar. This chapter is fundamental not just for your Class 11 understanding but forms the bedrock for many questions in competitive government exams. Pay close attention to the trends and, very importantly, the exceptions.

Chapter 3: Classification of Elements and Periodicity in Properties - Detailed Notes

1. Need for Classification & Historical Development:

• Why Classify? To organize the study of numerous elements and their compounds systematically, correlate their properties, and predict properties of undiscovered elements.
• Dobereiner's Triads (1829): Grouped elements in threes (triads) where the atomic weight of the middle element was roughly the arithmetic mean of the other two (e.g., Li, Na, K; Ca, Sr, Ba). Limited applicability.
• Newlands' Law of Octaves (1865): Arranged elements in increasing order of atomic weight and observed that every eighth element had properties similar to the first, like musical notes. Worked well only up to Calcium. Rejected noble gases.
• Mendeleev's Periodic Law (1869): "The properties of the elements are a periodic function of their atomic weights."
  • Mendeleev's Periodic Table: Arranged elements in periods (horizontal rows) and groups (vertical columns) based on increasing atomic weight.
  • Merits:
    • Systematic study of elements.
    • Correction of doubtful atomic weights (e.g., Be, In, U).
    • Prediction of new elements (eka-aluminium/Gallium, eka-silicon/Germanium) and their properties. Left gaps for them.
  • Demerits:
    • Anomalous pairs (Ar before K, Co before Ni, Te before I) - based on properties, not strictly atomic weight.
    • Position of isotopes (same element, different atomic weights) not defined.
    • Position of hydrogen was uncertain.
    • Lanthanoids and Actinoids placement.

2. Modern Periodic Law and the Present Form of the Periodic Table:

• Modern Periodic Law (Moseley, 1913): "The physical and chemical properties of the elements are periodic functions of their atomic numbers."
  • Atomic Number (Z) = number of protons (or electrons in a neutral atom). It's a more fundamental property than atomic mass.
• Present Form (Long Form) of the Periodic Table: Based on the Modern Periodic Law and electronic configuration.
  • Periods: 7 horizontal rows. The period number corresponds to the principal quantum number (n) of the outermost shell.
    • 1st Period (n=1): 2 elements (H, He) - Very Short
    • 2nd Period (n=2): 8 elements (Li to Ne) - Short
    • 3rd Period (n=3): 8 elements (Na to Ar) - Short
    • 4th Period (n=4): 18 elements (K to Kr) - Long
    • 5th Period (n=5): 18 elements (Rb to Xe) - Long
    • 6th Period (n=6): 32 elements (Cs to Rn) - Very Long (includes Lanthanoids)
    • 7th Period (n=7): Incomplete, currently 32 elements (Fr to Og) - Very Long (includes Actinoids)
  • Groups: 18 vertical columns. Elements in the same group have similar valence shell electronic configurations and hence similar chemical properties.
  • Blocks: Based on the subshell into which the last electron enters.
    • s-Block: Groups 1 (alkali metals) & 2 (alkaline earth metals). Last electron enters the ns subshell. General configuration: ns¹⁻². Highly reactive metals, form ionic compounds (except Li, Be).
    • p-Block: Groups 13 to 18. Last electron enters the np subshell. General configuration: ns²np¹⁻⁶. Includes metals, non-metals, and metalloids. Form mostly covalent compounds. Group 18 are Noble Gases (ns²np⁶, except He: 1s²). Group 17 are Halogens. Group 16 are Chalcogens.
    • d-Block: Groups 3 to 12. Last electron enters the (n-1)d subshell. General configuration: (n-1)d¹⁻¹⁰ ns⁰⁻². Known as Transition Elements. Exhibit variable oxidation states, form coloured ions, act as catalysts, often paramagnetic. Zn, Cd, Hg (Group 12, (n-1)d¹⁰ ns²) are generally not considered typical transition elements as their d-subshell is completely filled in the ground state and common oxidation state (+2).
    • f-Block: Placed separately at the bottom. Last electron enters the (n-2)f subshell. General configuration: (n-2)f¹⁻¹⁴ (n-1)d⁰⁻¹ ns². Known as Inner Transition Elements.
      • Lanthanoids: (n=6, 4f filling) Ce (Z=58) to Lu (Z=71). Follow Lanthanum (La, Z=57).
      • Actinoids: (n=7, 5f filling) Th (Z=90) to Lr (Z=103). Follow Actinium (Ac, Z=89). Mostly radioactive.

3. Nomenclature of Elements with Z > 100:

• Systematic IUPAC nomenclature using numerical roots for digits 0-9.
  • 0=nil(n), 1=un(u), 2=bi(b), 3=tri(t), 4=quad(q), 5=pent(p), 6=hex(h), 7=sept(s), 8=oct(o), 9=enn(e).
• Roots are assembled in order of digits, ending with '-ium'.
• Example: Z=104 -> un + nil + quad + ium = Unnilquadium (Unq)
• Example: Z=118 -> un + un + oct + ium = Ununoctium (Uuo) - Now named Oganesson (Og).

4. Periodic Trends in Properties:

(a) Atomic Radius: Half the internuclear distance between combined atoms.
* Types:
* Covalent Radius: Half the distance between nuclei of two covalently bonded atoms of the same element in a molecule.
* Metallic Radius: Half the internuclear distance separating metal cores in the metallic crystal.
* van der Waals Radius: Half the distance between the nuclei of two identical non-bonded isolated atoms or adjacent atoms in neighbouring molecules in the solid state. (Largest radius type). van der Waals > Metallic > Covalent.
* Trends:
* Across a Period (Left to Right): Decreases. Reason: Effective nuclear charge (Zeff) increases as electrons are added to the same shell, pulling the shell closer.
* Down a Group (Top to Bottom): Increases. Reason: Addition of a new electron shell (increase in principal quantum number 'n'), which outweighs the increase in nuclear charge. Screening effect also increases.
* Note: Noble gases have the largest atomic radii (van der Waals) in their respective periods but are usually compared separately.

(b) Ionic Radius: Radius of an ion in an ionic crystal.
* Cation: Smaller than the parent atom (loss of electrons, increased effective nuclear charge).
* Anion: Larger than the parent atom (gain of electrons, increased electron-electron repulsion, decreased effective nuclear charge).
* Isoelectronic Species: Ions/atoms having the same number of electrons (e.g., N³⁻, O²⁻, F⁻, Ne, Na⁺, Mg²⁺, Al³⁺). For isoelectronic species, radius decreases as atomic number (nuclear charge) increases. (Al³⁺ < Mg²⁺ < Na⁺ < Ne < F⁻ < O²⁻ < N³⁻).

(c) Ionization Enthalpy (IE) or Ionization Potential: Minimum energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state. (Unit: kJ mol⁻¹)
* A(g) + Energy → A⁺(g) + e⁻ (First IE, ΔᵢH₁)
* A⁺(g) + Energy → A²⁺(g) + e⁻ (Second IE, ΔᵢH₂)
* Successive IE: ΔᵢH₁ < ΔᵢH₂ < ΔᵢH₃ ... (It's harder to remove an electron from a positive ion).
* Factors Affecting IE:
* Atomic Size: IE decreases as atomic size increases.
* Nuclear Charge: IE increases as nuclear charge increases.
* Screening/Shielding Effect: Inner shell electrons shield valence electrons from the nucleus. IE decreases as screening effect increases.
* Penetration Effect: Order of penetration (closeness to nucleus): s > p > d > f. Electrons in more penetrating subshells are harder to remove (higher IE). E.g., IE of Be (1s²2s²) > B (1s²2s²2p¹).
* Electronic Configuration: Atoms with stable half-filled (p³, d⁵, f⁷) or fully-filled (p⁶, d¹⁰, f¹⁴) subshells have higher IE than expected. E.g., IE of N (1s²2s²2p³) > O (1s²2s²2p⁴); IE of Noble gases is highest in their periods.
* Trends:
* Across a Period: Generally increases. Reason: Increased Zeff and decreased atomic size. Exceptions: Be > B; N > O; Mg > Al; P > S.
* Down a Group: Generally decreases. Reason: Increased atomic size and screening effect, outweighing increased nuclear charge.

(d) Electron Gain Enthalpy (EGE, Δ<0xE2><0x82><0x91>H): Enthalpy change when an electron is added to an isolated gaseous atom in its ground state. (Unit: kJ mol⁻¹)
* X(g) + e⁻ → X⁻(g) ; Δ<0xE2><0x82><0x91>H
* Exothermic Process (Energy Released): Δ<0xE2><0x82><0x91>H is negative (favourable electron gain). More negative value means higher tendency to accept an electron.
* Endothermic Process (Energy Absorbed): Δ<0xE2><0x82><0x91>H is positive (unfavourable electron gain).
* Factors Affecting EGE: Similar to IE (Atomic size, Nuclear Charge, Electronic Configuration).
* Trends:
* Across a Period: Generally becomes more negative (more exothermic). Reason: Increased Zeff and decreased atomic size favour electron addition.
* Down a Group: Generally becomes less negative (less exothermic). Reason: Increased atomic size and screening effect make electron addition less favourable.
* Exceptions/Important Points:
* Halogens (Group 17): Have the most negative EGE values (highest tendency to gain an electron). Order: Cl > F > Br > I. (EGE of F is less negative than Cl due to small size and high electron density in F, leading to strong interelectronic repulsions). This F < Cl exception is very important.
* Noble Gases (Group 18): Have large positive EGE values (highly unfavourable) due to stable ns²np⁶ configuration. Electron has to enter the next higher shell (n+1).
* Be, Mg, Ca (Group 2): Have positive EGE values. Electron needs to enter the higher energy np subshell (ns² configuration is stable).
* N, P (Group 15): Have near zero or slightly positive/less negative EGE than Group 14/16 elements due to stable half-filled np³ configuration. P has a more negative EGE than N (similar reason as F vs Cl).
* Successive EGE: Addition of a second electron (e.g., O⁻(g) + e⁻ → O²⁻(g)) is always endothermic (positive Δ<0xE2><0x82><0x91>H₂) due to repulsion between the anion and the incoming electron.

(e) Electronegativity (EN): Qualitative measure of the ability of an atom in a chemical compound to attract the shared pair of electrons towards itself. (Dimensionless property).
* Scales: Pauling scale, Mulliken scale (EN = (IE + EGE)/2). Pauling scale is most common. F is assigned EN=4.0.
* Factors Affecting EN: Similar to IE (Atomic size, Nuclear charge, Zeff). Also depends on the state of hybridization (EN order: sp > sp² > sp³).
* Trends:
* Across a Period: Increases. Reason: Increased Zeff.
* Down a Group: Decreases. Reason: Increased atomic size.
* Key EN Values (Pauling): F(4.0) > O(3.5) > N(3.0) ≈ Cl(3.0) > Br(2.8) > I(2.5) ≈ S(2.5) ≈ C(2.5) > H(2.1) ≈ P(2.1). Metals generally have low EN (< 2.0), Non-metals have high EN (> 2.0). Cs is the least electronegative (most electropositive), F is the most electronegative.

(f) Valency and Oxidation State:
* Valency: Usually equals the number of valence electrons or (8 - number of valence electrons). For s-block and p-block elements, the valency with respect to oxygen is usually the group number, and with respect to hydrogen is group number (Groups 1, 2, 13, 14) or (18 - group number) (Groups 15, 16, 17).
* Oxidation State: Apparent charge on an atom in a compound. An element can exhibit multiple oxidation states (especially p-block, d-block).
* Trends: Oxidation states vary periodically. The highest oxidation state for representative elements (s & p block) is often the group number (e.g., +7 for Cl in Cl₂O₇).

(g) Metallic and Non-metallic Character:
* Metallic Character: Tendency to lose electrons (low IE, low EN). Increases down a group, decreases across a period.
* Non-metallic Character: Tendency to gain electrons (high IE, high EN). Decreases down a group, increases across a period.
* Metalloids (Semi-metals): Elements showing properties intermediate between metals and non-metals (e.g., Si, Ge, As, Sb, Te). Lie along the diagonal border between metals and non-metals.

(h) Nature of Oxides:
* Across a Period: Basic → Amphoteric → Acidic.
* Na₂O (Strongly Basic), MgO (Basic), Al₂O₃ (Amphoteric), SiO₂ (Acidic), P₄O₁₀ (Strongly Acidic), SO₃ (Strongly Acidic), Cl₂O₇ (Very Strongly Acidic).
* Down a Group: Basic character of oxides generally increases (metallic character increases). E.g., BeO (Amphoteric), MgO (Basic), CaO (Basic), SrO (Basic), BaO (Strongly Basic). For non-metal oxides, acidic character tends to decrease down the group.
* Amphoteric Oxides: React with both acids and bases (e.g., Al₂O₃, ZnO, BeO, Ga₂O₃, SnO, PbO).
* Neutral Oxides: Do not react with acids or bases (e.g., CO, NO, N₂O).

5. Anomalous Properties of Second Period Elements:

• Elements of the 2nd period (Li, Be, B, C, N, O, F) show significant differences from other elements in their respective groups.
• Reasons:
  • Small atomic size.
  • High Ionization Enthalpy and High Electronegativity.
  • Absence of vacant d-orbitals in their valence shell (maximum covalency is 4).
• Examples: Li forms covalent compounds unlike other alkali metals. Be forms covalent compounds (unlike Mg, Ca...). Boron is a metalloid, forms acidic oxide (unlike Al - amphoteric). Carbon forms pπ-pπ multiple bonds (C=C, C≡C, C=O) readily, Si does not. N₂ is a diatomic gas with triple bond, P₄ is a solid. O₂ is a diatomic gas, S₈ is a solid. HF is a liquid (H-bonding), HCl is a gas.

6. Diagonal Relationship:

• Similarity in properties between elements of the second period and elements diagonally placed in the third period.
  • Li ≈ Mg
  • Be ≈ Al
  • B ≈ Si
• Reason: Similar ionic size and/or charge/radius ratio (ionic potential, φ = charge/radius). They have similar electronegativity values.
• Examples: Li and Mg form nitrides, carbonates decompose easily, chlorides are deliquescent and soluble in ethanol. Be and Al form covalent compounds, oxides are amphoteric, chlorides are Lewis acids and exist as dimers. B and Si are semiconductors, form acidic oxides, form covalent hydrides and halides.

Multiple Choice Questions (MCQs)

1. Which of the following pairs of atomic numbers represents elements belonging to the s-block?
   (a) 7, 15
   (b) 12, 20
   (c) 9, 17
   (d) 3, 13

2. The element with atomic number 117 (Tennessine, Ts) belongs to which group of the periodic table?
   (a) Group 1
   (b) Group 2
   (c) Group 17
   (d) Group 18

3. Consider the isoelectronic species: Na⁺, Mg²⁺, F⁻, O²⁻. The correct order of increasing ionic radii is:
   (a) O²⁻ < F⁻ < Na⁺ < Mg²⁺
   (b) Mg²⁺ < Na⁺ < F⁻ < O²⁻
   (c) Mg²⁺ < O²⁻ < Na⁺ < F⁻
   (d) F⁻ < O²⁻ < Na⁺ < Mg²⁺

4. Which of the following elements has the highest first ionization enthalpy?
   (a) Boron
   (b) Carbon
   (c) Nitrogen
   (d) Oxygen

5. Which element exhibits the most negative electron gain enthalpy?
   (a) Fluorine
   (b) Chlorine
   (c) Oxygen
   (d) Sulphur

6. The correct order of electronegativity among the following elements is:
   (a) F > N > O > C
   (b) F > O > C > N
   (c) O > F > N > C
   (d) F > O > N > C

7. Which of the following oxides is amphoteric in nature?
   (a) CaO
   (b) CO₂
   (c) SiO₂
   (d) Al₂O₃

8. The anomalous behaviour of Lithium compared to other alkali metals is primarily due to its:
   (a) High ionization enthalpy
   (b) Small size
   (c) High electronegativity
   (d) All of the above

9. The IUPAC name for the element with atomic number 109 is:
   (a) Unnilennium
   (b) Unnilbium
   (c) Ununennium
   (d) Unnilennium

10. Which property generally decreases down a group in the periodic table?
    (a) Atomic Radius
    (b) Ionization Enthalpy
    (c) Metallic Character
    (d) Basic nature of oxides

Answer Key for MCQs:

1. (b) [Mg (12) and Ca (20) are Group 2 elements]
2. (c) [Element 117 is below Fluorine/Chlorine, hence Group 17 (Halogens)]
3. (b) [For isoelectronic species, radius decreases as nuclear charge increases]
4. (c) [Nitrogen has a stable half-filled p³ configuration, making its IE higher than O]
5. (b) [Chlorine has a more negative EGE than Fluorine due to less interelectronic repulsion in the larger 3p subshell compared to 2p]
6. (d) [Standard EN trend: F > O > N > C]
7. (d) [Al₂O₃ reacts with both acids and bases]
8. (d) [Small size, high IE, and high EN all contribute to Li's anomalous behaviour]
9. (d) [1=Un, 0=nil, 9=enn => Unnilennium (Une)] Note: Official name is Meitnerium (Mt). The question asks for IUPAC systematic name.
10. (b) [Ionization Enthalpy decreases down a group due to increased size and shielding]

Study these notes thoroughly, focusing on the trends, the reasons behind them, and especially the exceptions like those in IE and EGE. Understanding the structure of the periodic table based on electronic configuration is crucial. Good luck with your preparation!